Atomic Structure PDF

Atomic Structure Organic Chemistry Organic Chemistry Organic chemistry is the study of carbon compounds. But why is carbon special? Why, of the more than 50 million presently known chemical compounds, do most of them contain carbon? Carbon can share four valence electrons and form four strong covale...

Atomic Structure Organic Chemistry Organic Chemistry Organic chemistry is the study of carbon compounds. But why is carbon special? Why, of the more than 50 million presently known chemical compounds, do most of them contain carbon? Carbon can share four valence electrons and form four strong covalent bonds. Organic Chemistry Furthermore, carbon atoms can bond to one another, forming long chains and rings. Carbon, alone of all elements, is able to form an immense diversity of compounds, from the simple methane, with one carbon atom, to the complex DNA, which can have more than 100 million carbons. Organic Chemistry Of course, not all carbon compounds are derived from living organisms. Modern chemists have developed a remarkably sophisticated ability to design and synthesize new organic compounds in the laboratory— medicines, dyes, polymers, and a host of other substances Atomic Structure Atomic Structure The Nucleus Atom consists of a dense, positively charged nucleus surrounded at a relatively large distance by negatively charged electrons. The nucleus consists of subatomic particles called protons, which are positively charged, and neutrons, which are electrically neutral. Because an atom is neutral overall, the number of positive protons in the nucleus and the number of negative electrons surrounding the nucleus are the same. Atomic Structure Orbitals In atomic theory and quantum mechanics an atomic orbital is a mathematical function describing the location and wave-like behavior of an electron in an atom. This function can be used to calculate the probability of finding any electron of an atom in any specific region around the atoms nucleus. What do orbitals look like? There are four different kinds of orbitals, denoted s, p, d, and f, each with a different shape. Of the four, we’ll be concerned primarily with s and p orbitals because these are the most common in organic and biological chemistry. Atomic Structure The orbitals in an atom are organized into different electron shells, centered around the nucleus and having successively larger size and energy. Different shells contain different numbers and kinds of orbitals, and each orbital within a shell can be occupied by two electrons. The first shell contains only a single s orbital, denoted 1s, and thus holds only 2 electrons. Atomic Structure The second shell contains one 2s orbital and three 2p orbitals and thus holds a total of 8 electrons. The third shell contains a 3s orbital, three 3p orbitals, and five 3d orbitals, for a total capacity of 18 electrons. These orbital groupings and their energy levels are shown in Figure. Atomic Structure Electron Configurations The lowest-energy arrangement, or ground-state electron configuration, of an atom is a listing of the orbitals occupied by its electrons. We can predict this arrangement by following three rules. Rule 1 The lowest-energy orbitals fill up first, according to the order 1s - 2s - 2p - 3s - 3p - 4s - 3d Atomic Structure Rule 2 Electrons act in some ways as if they were spinning around an axis, somewhat like how the earth spins. This spin can have two orientations, denoted as up ( ) and down ( ). Only two electrons can occupy an orbital, and they must be of opposite spin, a statement called the Pauli exclusion principle. Atomic Structure Rule 3 If two or more empty orbitals of equal energy are available, one electron occupies each with spins parallel until all orbitals are half-full, a statement called Hund’s rule. Some examples of how these rules apply are shown in this Table Shapes of Atomic Orbitals The Ionic Bond The ionic bond Chemical bonds Are the forces that hold atoms together in a molecule. 1. The ionic bond Results from transfer of electrons, as, for example, in the formation of lithium fluoride. A lithium atom has two electrons in its inner shell and one electron in its outer or valence shell ; the loss of one electron would leave lithium with a full outer shell of two electrons. The ionic bond A fluorine atom has two electrons in its inner shell and seven electrons in its valence shell; the gain of one electron would give fluorine a full outer shell of eight. Lithium fluoride is formed by the transfer of one electron from lithium to fluorine; lithium now bears a positive charge and fluorine bears a negative charge. The electrostatic attraction between the oppositely charged ions is called an ionic bond. The ionic bond Ionic Bonding Ionic Bond is the chemical bond that forms when valence electron(s) is completely transferred between atoms and that bond generates two oppositely charged ions. The Covalent Bond The covalent bond 2- The covalent bond Results from sharing of electrons, as, for example, in the formation of the hydrogen molecule. Each hydrogen atom has a single electron; by sharing a pair of electrons, both hydrogen can complete their shells of two. Two fluorine atoms, each with seven electrons in the valence shell, can complete their octets by sharing a pair of electrons. In a similar way we can visualize the formation of HF, H2O, NH3, CH4, and CF4. The covalent bond The covalent bond The covalent bond is typical of the compounds of carbon; it is the bond of chief importance in the study of organic chemistry. The covalent bond Why, though, do atoms bond together, and how can bonds be described electronically? The why question is relatively easy to answer: atoms bond together because the compound that results is more stable and lower in energy than the separate atoms. Energy—usually as heat—always flows out of the chemical system when a bond forms. Conversely, energy must be put into the chemical system to break a bond. Making bonds always releases energy, and breaking bonds always absorbs energy. The covalent bond The how question is more difficult. To answer it, we need to know more about the electronic properties of atoms. We know through observation that eight electrons (an electron octet) in an atom’s outermost shell, or valence shell, impart special stability to the noble gas elements in group 8 A of the periodic table: Ne (2 + 8); Ar (2 + 8 + 8); Kr (2 + 8 + 18 +8). We also know that the chemistry of main-group elements is governed by their tendency to take on the electron configuration of the nearest noble gas. The covalent bond The neutral collection of atoms held together by covalent bonds is called a molecule. A simple way of indicating the covalent bonds in molecules is to use what are called Lewis structures, or electron-dot structures, in which the valence shell electrons of an atom are represented as dots. Thus, hydrogen has one dot representing its 1s electron, carbon has four dots (2s2 2p2), oxygen has six dots (2s2 2p4), and so on. A stable molecule results whenever a noble-gas configuration is achieved for all the atoms—eight dots (an octet) for main-group atoms or two dots for hydrogen. The covalent bond Simpler still is the use of Kekulé structures, or line-bond structures, in which a twoelectron covalent bond is indicated as a line drawn between atoms. The covalent bond The number of covalent bonds an atom forms depends on how many additional valence electrons it needs to reach a noble-gas configuration. Hydrogen has one valence electron (1s) and needs one more to reach the helium configuration (1s2), so it forms one bond. Carbon has four valence electrons (2s2 2p2) and needs four more to reach the neon configuration (2s2 2p6), so it forms four bonds. 3 Sp Hybrid Orbitals Sp3 Hybrid Orbitals and the Structure of Methane (CH4) Carbon has four valence electrons (2s2 2p2) and forms four bonds. Because carbon uses two kinds of orbitals for bonding, 2s and 2p, we might expect methane to have two kinds of C-H bonds. In fact, though, all four C-H bonds in methane are identical and are spatially oriented toward the corners of a regular tetrahedron. 3 Sp Hybrid Orbitals How can we explain this? An answer was provided by Linus Pauling, who showed mathematically how an s orbital and three p orbitals on an atom can combine, or hybridize, to form four equivalent atomic orbitals with tetrahedral orientation. These tetrahedral oriented orbitals are called sp3 hybrid orbitals. Note that the superscript 3 in the name sp3 tells how many of each type of atomic orbital combine to form the hybrid, not how many electrons occupy it. 3 Sp Hybrid Orbitals When an s orbital hybridizes with three p orbitals, the resultant sp3 hybrid orbitals are unsymmetrical about the nucleus. One of the two lobes is larger than the other and can therefore overlap more effectively with an orbital from another atom to form a bond. As a result, sp3 hybrid orbitals form stronger bonds than do unhybridized s or p orbitals. 3 Sp Hybrid Orbitals The angle formed by each H - C - H is 109.5°, called tetrahedral angle. 2 Sp Hybrid Orbitals sp2 Hybrid Orbitals and the Structure of Ethylene The bonds we’ve seen in methane called single bonds because they result from the sharing of one electron pair between bonded atoms. However, that carbon atom can also form double bonds by sharing two electron pairs between atoms or triple bonds by sharing three electron pairs. 2 Sp Hybrid Orbitals Ethylene, for instance, has the structure H2C=CH2 and contains a carbon–carbon double bond, while acetylene has the structure HC ≡ CH and contains a carbon–carbon triple bond. 2 Sp Hybrid Orbitals Imagine instead that the 2s orbital combines with only two of the three available 2p orbitals. Three sp2 hybrid orbitals result and one 2p orbital remains unchanged. Like sp3 hybrids, sp2 hybrid orbitals are unsymmetrical about the nucleus and are strongly oriented in a specific direction so they can form strong bonds. The three sp2 orbitals lie in a plane at angles of 120° to one another, with the remaining p orbital perpendicular to the sp2 plane 2 Sp Hybrid Orbitals When two carbons with sp2 hybridization approach each other, they form a strong σ bond by sp2–sp2 head-on overlap. At the same time, the unhybridized p orbitals interact by sideways overlap to form what is called a pi (π) bond. The combination of an sp2–sp2 σ bond and a 2p–2p π bond results in the sharing of four electrons and the formation of a carbon–carbon double bond. Sp Hybrid Orbitals sp Hybrid Orbitals and the Structure of Acetylene In addition to forming single and double bonds by sharing two and four electrons, respectively, carbon also can form a triple bond by sharing six electrons. To account for the triple bond in a molecule such as acetylene, H-C ≡ C-H, we need a third kind of hybrid orbital, an sp hybrid.. 2 Sp Hybrid Orbitals Imagine that, instead of combining with two or three p orbitals, a carbon 2s orbital hybridizes with only a single p orbital. Two sp hybrid orbitals result and two p orbitals remain unchanged. The two sp orbitals are oriented 180° apart on the x-axis. 2 Sp Hybrid Orbitals When two sp-hybridized carbon atoms approach each other, sp hybrid orbitals on each carbon overlap head-on to form a strong sp–sp σ bond. At the same time, the pz orbitals from each carbon form a pz–pz π bond by sideways overlap, and the py orbitals overlap similarly to form a py–py π bond.